Section 1

Chemical bond- a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

Covalent bonding- Results from the sharing of electron pairs between two atoms.

Ionic bonds- Chemical bonds that result from the electrical attraction between cations and anions. In these bonds, atoms completely give or lose electrons to other atoms. [Difference in electronegativity is greater than or equal to 2.0]

Polar Covalent Bond

Polar-covalent bonds- A covalent bond in which the bonded atoms have an uneven distribution of charge. [Difference in electronegativity is 0.5-1.9]

Non-polar covalent- A covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. [Difference in electronegativity is 0.0- 0.4]

Practice Problems

1. What is the electronegativity difference if Bromine and Hydrogen form a bond? What kind of bond would it be?
2. Name two elements that would form a non-polar covalent bond.
3. If Cesium forms a bond that has an electronegativity difference of 2.8, what element is it bonded to?

Section 2

Molecular compound- A chemical compound whose simplest units are molecules.

• The composition of a compound is given by its chemical formula, which indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

Electron-Dot Notation- An electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element’s symbol.

Molecule

Molecule- A neutral group of atoms that are held together by covalent bonds.

• They are able to exist on its own
• They can consist of two or more atoms of the same element, or of different elements

Formation of Covalent Bonds:

• Most atoms have lower potential energy when they are bonded to other atoms than they have as independent particles.
• As atoms move towards each other, the nuclei and electrons are attracted to each other. The two nuclei repel each other and the two electrons repel each other. Attraction between particles corresponds to a decrease in potential energy of the atoms, while repulsion corresponds to an increase in potential energy.
• When the atoms first notice each other’s presence, the electron-proton attraction is stronger than the electron-electron and proton-proton repulsions, so the atoms are drawn together and their potential energy is lowered. The total energy continues to be lowered as they approach each other until a distance is reached at which the repulsion between the like charges equals the attraction of the opposite charges. At this point, a stable molecule forms.

Characteristics of Covalent Bonds:

• The bond length is the average distance between the nuclei of two covalently bonded atoms
• Bond energy is the energy required to break a chemical bond and form neutral isolated atoms. This is measured in kJ/mol (indicates the energy required to break one mole of bonds in isolated molecules)
• Vary with the types of atoms being combined
• While forming a covalent bond, atoms release energy as they change from isolated individual atoms to parts of a molecule (the amount of potential energy released equals the difference between the potential energy at the zero level and that of bonded atoms). The same amount of energy must be added to separate the bonded atoms.

The Octet Rule:

• Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
  • Exception: some elements can be surrounded by less than 8 electrons. Examples of this would be Hydrogen and Boron, which have only two and three valence electrons, respectively.
  • Exception: some elements can be surrounded by more than 8 electrons when they combine with the highly electronegative elements fluorine, oxygen and chlorine. In these bonds, electrons in d orbitals as well as in s and p orbitals are involved. Examples of this are PF5 and SF6.

Lewis Structures:

• Electron-dot notation is used to represent molecules instead of individual atoms.
• Atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons.
  • Example: H:H is called a single bond (a covalent bond in which one pair of electrons is shared between two atoms)
• A lone pair is a pair of electrons that belongs to one atom and does not bond at all with the other atom
• A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule.

Resonance Structures (or resonance hybrids):

• Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.
• To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures.

Section 3

Ionic Bonding and Ionic Compounds Overview

The ratio of ions in a formula unit depends on the charges of the ions combine. Example: Fluoride = Two Fluoride anions each with negative charges. This must be balanced by two positively charged ions which Ca2+: CaF2

Structure of Sodium Chloride illustrating atom positioning

In ionic bonding, the bonding ions work to reduce their potential energy by arranging themselves in a certain way known as "crystal lattice." (Example displayed at right). These forces of repulsion between atoms causes the charged ions to extend over large distances in a three-dimensional plane ultimately causing the potential energy to decrease. The size of the three-dimensional expansion depends entirely on the sizes, charges, and amount of ions.

The boiling point, melting point, and hardness of a compound depends on the attraction of its elements. Forces of attraction between individual molecules are very weak causing them to have low melting points, this explains why many molecular compounds are in gas form in room temperature.

Key facts concerning comparing ionic and molecular compounds:

• In the solid state, ions cannot move. Consequently, they do not conduct electricity
• In the melted state, however, the ions can move about freely and become good conductors of electricity
• Most ionic compounds dissolve in water (ions seperate and are surrounded by water molecules).

Polyatomic ions are combined with ions of opposite charge to form "ionic compounds." When this occurs, there are an excess of electrons and a shortage of protons. These combinations result in compounds that have properties of both molecular and ionic compounds.

Section 4

Overview:

In most metallic atoms, the highest orbital is occupied by very few, if any, electrons. Many of them possess d orbitals that are entirely empty. Because of this lack of fulfillment, these vacant outer orbitals overlap with other orbitals, allowing the outer electrons to move about freely, creating a "sea of electrons".

Metallic bonding- The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons

Section 5

VSEPR Theory

• VSEPR = Valence Shell Electrons-Pair Repulsion
• VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.
• Shared pairs will be as far away from each other as possible

Practice Problems

Complete on a seperate sheet of paper.

1. What kind of spatial geometry do the compounds H20, CH4, and KCl form? Draw them.
2. Draw the electron dot diagrams for As, Ar, and Li.